Download Light and quantized Energy Section 1

Quiz Review
Light and quantized Energy
Section 1
Electrons in Atoms
Chapter 5
Unanswered Questions

What have we learned?

Subatomic particles exist

The nucleus of an atom contains all of the atoms
positive charge (Protons) and mass (Neutrons +
Protons)

Atoms are surrounded by fast, moving electrons
Unanswered Questions

But there were some unanswered
questions:

How the atoms electrons were arranged in space

Why the electrons were not pulled into the
positively charged nucleus.

Why there were differences and similarities in the
chemical behavior among various elements
Unanswered Questions
 Unraveling of the puzzle:

Scientists observed that certain elements emitted
visible light when heated in a flame.

Analysis of this light revealed that an elements
chemical behavior is related to the arrangement
of the electrons in its atoms.

To understand this we have to first understand
the wave nature of light
The Wave Nature of Light
 Electromagnetic Radiation: Is a form of
energy that exhibits wave like behavior as it
travels through space

Visible light is a type of electromagnetic radiation

Other examples include those from:
 Microwaves
 X-rays
 Signals
home
sent from the Radio & TV station to your
The Wave Nature of Light
The Wave Nature of Light


Characteristics of Waves
Wavelength (λ): the shortest distance between
equivalent points on a continuous wave.
 Usually expressed in meters (m)
Frequency (ν): the number of waves that pass a given
point per second.
 Expressed as:

Hertz (Hz) = One wave per second

Waves = Waves/ second (1/s) or (s-1)

Example: 652 Hz = 652 waves/second = 652/s = 652 s-1
The Wave Nature of Light


Characteristics of Waves
Amplitude: Wave’s height from the origin to a crest or
origin to trough
Electromagnetic Wave Relationship:
 Speed of Electromagnetic waves = 3.00x108 m/s
 Speed of light (c) = 3.00x108 m/s

Equation: c = λv where:
c
= 3.00x108 m/s
 λ = wave length
 ν = frequency
The Wave Nature of Light
Waves can have different wavelengths & frequencies
Wavelengths and frequency are inversely related
Electromagnetic Spectrum
Practice Problems 1—4 page 140
Practice Problems 1—4 page 140
Practice Problems 1—4 page 140
Practice Problems 1—4 page 140
The Particle Nature of Light

The wave nature of light does not explain certain
phenomenon



Light emission of certain metals
Photoelectric effect
Studied by German physicist Max Planck (1858-1947) in
1900
The Particle Nature of Light

Planck’s conclusions:



Matter can gain or lose energy only in small, specific amounts
called quanta.
Quantum: minimum amount of energy that can be gained or
lost by an atom
Planck’s constant (h) = 6.626 × 10-34 J ● s
The Photoelectric Effect

Photoelectric effect: electrons are emitted from the
surface of a metal when light of a certain frequency
shines on the metal
The Photoelectric Effect



Albert Einstein proposed in 1905 that light has a dual
nature.
“A beam of light has wavelike and particle-like
properties.”
Light particle: photon

a particle of electromagnetic radiation with no mass that
carries a quantum of energy
The Photoelectric Effect
Practice Problems 5—7 page 143
Atomic Emission Spectrum
Atomic Emission Spectrum
 Atomic
emission spectrum of an
element:

the set of frequencies of the electromagnetic
waves emitted by the atoms of the element
Atomic Emission Spectrum Cont’d

Unique for each element and can be used to identify
an element or determine whether an element is part
of an unknown compound.
Atomic Emission Spectrum Cont’d

The atomic emission spectrum is not continuous
(made up of certain frequencies of light)
Quantum Theory and the Atom
Section 2
Electrons in Atoms
Chapter 5
Bohr’s model of the atom

Einstein’s theory of light’s dual nature accounted for
several unexplainable phenomena (photoelectric effect)
but not why atomic emission spectra of elements were
discontinuous rather continuous.

In 1913, Niels Bohr, a Danish physicist working in
Rutherford’s laboratory, proposed a quantum model for
the hydrogen atom that seemed to answer this question.
Bohr’s model of the atom

Elements have certain allowable energy
states
Ground State: Lowest energy state of an atom
 Excited Sate: The state of an atom when it
gains energy

Bohr’s model of the atom Cont’d


An atoms energy state is related to the location of
the electron.
Electron moves around the nucleus only in certain
allowed circular orbits
The smaller the electrons orbit
the lower the atoms energy
state or energy level
The larger the electrons orbit
the higher the atoms energy
state or energy level
Bohr’s model of the atom Cont’d

The energy state of an electron is related to its quantum
number

The quantum number is number of the orbit in an atom

Ex.Hydrogen has one electron which exists in
the ground state or quantum # of n= 1

When energy is absorbed the electron moves to a
higher energy orbit e.g. n=2 (the excited state)

The electron of the atom can return from:
The excited state n=2 to its
 Ground state n=1 by emitting (releasing)
energy as a photon


This energy difference can be calculated using the
formula

Δ E = E higher-energy orbit – E lower-energy orbit = E photon = hv
The Quantum Mechanical Model of
the Atom

Bohr’s model was flawed: Electrons do not move in a
circular orbit

Heisenberg uncertainty principle: It is fundamentally
impossible to know precisely both the velocity and
position of a particle at the same time
The Quantum Mechanical Model of the
Atom

The quantum mechanical model makes no attempt
to describe the electrons path around nucleus

Boundary encompasses the 95% probability that the
electron is located there at any given time
The Quantum Mechanical Model of
the Atom

Uses 4 numbers to “address” an electron in an atom.
We will only work with the principal quantum #
 Principal quantum number, n
Energy Sublevels is equal to the number of
the principal quantum number
The Quantum Mechanical Model of
the Atom
Shape of Orbitals
 Sublevels
are labelled s, p, d or f depending
on the shape of the orbital
os- Spherical
op – Dumbbell shaped
od – Same shape
different planes
of – Same shape
different planes
The Quantum Mechanical Model of
the Atom
s
Orbitals

Have only 1 orbital

Increase in size with increasing quantum number
The Quantum Mechanical Model of
the Atom
p
Orbitals

Have 3 orbitals

Oriented along three axes: x, y & z
The Quantum Mechanical Model of
the Atom
d
Orbitals

Have 5orbitals
Oriented along different planes: xy, xz, yz, x2-y2

Z2 – shaped differently oriented differently

The Quantum Mechanical Model of
the Atom
f


Orbitals
Have 7 orbitals
Complex multilobed shapes
The Quantum Mechanical Model of
the Atom
Shape of Orbitals
 Each
orbital contains at most 2 electrons
Principal
Quantum #
Sublevels (type
of Orbitals)
# of orbitals
related to sublevel
1s
1
2s
2p
1
3
n=3
3s
3p
3d
1
3
5
n=4
3s
3p
3d
3f
1
3
5
7
n=1
n=2
The Quantum Mechanical Model of
the Atom

Uses four numbers to “address” an electron in an
atom
 Principal quantum number, n
Energy Sublevels is equal to the number of
the principal quantum number
Angular momentum quantum number, l
 Magnetic quantum number, m
 Spin quantum number, + ½ or – ½

Electron Configuration
Section 3
Electrons in Atoms
Chapter 5
Ground state electron configuration

Electron Configuration: Arrangement of electrons
in an atom

The Aufbau principle: Each electron occupies the lowest
energy orbital available because they are more stable at this
level.

Hund’s rule: Single electrons with the same spin must occupy
each equal-energy orbital before additional electrons with
opposite spins can occupy the same energy level orbitals
Ground state electron configuration

The Pauli exclusion principle: A maximum of two
electrons can occupy a single orbital, but only if the electrons
have opposite spins
Using the Aufbau principle: Each electron occupies the
lowest energy orbital available because they are more stable
at this level.
Order of filling orbitals
Order of filling orbitals
The Pauli exclusion principle: A maximum of
two electrons can occupy a single orbital, but only if
the electrons have opposite spins
Hunts Rule: A maximum of two electrons can occupy a single
orbital, but only if the electrons have opposite spins
Ground state electron configuration
Ground state electron configuration
Writing the electron configuration
Neon
Electron Configuration: 1s2 2s2 2p6
Ground state electron configuration

Only works perfectly for all elements up to
and including Vanadium, atomic number 23
Practice Problem 21-25 p 160
Ground state electron configuration

Noble Gas notation: A method of representing
electron configurations of noble gases.
Ground state electron configuration

Noble Gases:
o
o

Have eight electrons in their outermost orbital
Usually unstable
Noble Gas Notation:
o
Uses bracketed symbols
o
For example: Helium (He)
Ground state electron configuration
o
For example:
o
For example:
Ground state electron configuration
o
For example:
o
For example:
Ground state electron configuration

Valence electrons
o
Defined as atom’s outermost orbitals
o
Atoms in the outermost orbitals are associated with the
atom’s highest principal energy level
o
For example:
Ground state electron configuration

Electron-dot structure
o
Devised by American Chemist G.N. Lewis
o
Aka Lewis Dot structures
o
Used by chemists to represent Valence Electrons
o
Writing the Lewis Dot Structure:
 Element’s Symbol: Represents the atomic nucleus and
inner level electrons.
Ground state electron configuration

o
Dots:
 Represent Valence electrons
 Placed one at a time around the four sides of the
symbol and then paired up until all are shown
Example:
Practice Problems
Pg. 162 #26-28
Practice Problems 26-28 p 162
Practice Problems 26-28 p 162
Practice Problems 26-28 p 162
Practice Problems 26-28 p 162
Practice Problems 26-28 p 162